Atoms bond to achieve a more stable electron configuration. A covalent bond is characterized by the sharing of valence electrons between two atoms. This contrasts with an ionic bond, which involves the complete transfer of electrons and the formation of charged ions. Identifying a covalent bond requires looking closely at the atoms involved.
The Quick Check: Location on the Periodic Table
A straightforward method for predicting a covalent bond involves locating the elements on the periodic table. Covalent bonds typically form when two nonmetal atoms combine. Nonmetals are found predominantly on the upper-right side of the periodic table, with the exception of hydrogen, which behaves as a nonmetal in most bonding scenarios.
When two atoms from this nonmetal region bond, such as carbon and oxygen in carbon dioxide (\(\text{CO}_2\)) or hydrogen and oxygen in water (\(\text{H}_2\text{O}\)), the resulting bond is almost always covalent. This provides a quick rule of thumb because nonmetal atoms possess similar electron-attracting tendencies. Neither atom is strong enough to completely remove an electron from the other, making electron sharing the most favorable outcome.
Conversely, the combination of a nonmetal with a metal suggests an ionic bond, where the metal transfers its electron to the nonmetal. While this periodic table location is a reliable first filter, it is a qualitative prediction. For definitive identification, a quantitative approach is necessary.
The Definitive Test: Calculating Electronegativity Difference
The most accurate way to identify a covalent bond is by calculating the difference in electronegativity (\(\Delta\text{EN}\)) between the two bonded atoms. Electronegativity is an intrinsic property that quantifies an atom’s ability to attract a shared pair of electrons toward itself within a chemical bond. The values for elements are typically measured on the Pauling scale, where fluorine is the most electronegative element with a value of 4.0.
The absolute difference between the electronegativity values of the two bonding atoms reveals the extent of electron sharing versus electron transfer. If the \(\Delta\text{EN}\) is very small, the sharing is relatively equal, confirming a covalent bond. As the difference increases, the sharing becomes unequal, and the bond gains more ionic character.
A bond is classified as covalent if the \(\Delta\text{EN}\) between the two atoms is less than 1.7. If the \(\Delta\text{EN}\) is less than 0.5, the bond is considered nonpolar covalent, indicating nearly equal sharing.
If the \(\Delta\text{EN}\) falls between 0.5 and the 1.7 threshold, the bond is classified as polar covalent. When the difference exceeds 1.7, the electron attraction is so uneven that the electron is considered essentially transferred, which defines an ionic bond. For example, a bond between Carbon (EN \(\approx 2.5\)) and Hydrogen (EN \(\approx 2.1\)) yields a \(\Delta\text{EN}\) of 0.4, placing it firmly in the nonpolar covalent range.
Refining the Identification: Polar Versus Nonpolar Bonds
Covalent bonds exist on a spectrum of sharing, refined by distinguishing between polar and nonpolar types. A nonpolar covalent bond occurs when the electron pair is shared equally, resulting from either two identical atoms (like in \(\text{O}_2\)) or two different atoms with a \(\Delta\text{EN}\) less than 0.5. In these cases, the electron density is distributed symmetrically across the bond.
A polar covalent bond forms when the \(\Delta\text{EN}\) is large enough to cause unequal sharing, between 0.5 and 1.7. The shared electrons are pulled more strongly toward the atom with the higher electronegativity. This unequal pull creates a partial negative charge (\(\delta^-\)) on the more electronegative atom and a partial positive charge (\(\delta^+\)) on the less electronegative atom.
This separation of charge within the bond is known as a bond dipole. While still sharing electrons, the bond behaves like a tiny magnet with distinct positive and negative ends. The degree of this polarity significantly influences the properties of the resulting molecule, affecting its melting point, solubility, and how it interacts with other polar substances like water.
A molecule can contain polar bonds yet be nonpolar overall if the dipoles are arranged symmetrically and cancel each other out, such as in carbon dioxide (\(\text{CO}_2\)). Understanding the precise degree of polarity, determined by the \(\Delta\text{EN}\), is the final step in accurately identifying and characterizing a covalent bond.